Download Naming inorganic compounds Ionic equations

Naming inorganic
compounds
Ionic equations
MUDr. Jan Pláteník, PhD
Atom
• Smallest particle of a pure element having its
chemical properties
• Positively charged nucleus (protons, neutrons)
• Negatively charged electron shell:
– Electron is wave/particle
– Behavior of electron described by quantum mechanics
(... wave function, quantum numbers)
– Orbital: space area within the atom shell where
occurrence of an electron or pair of electrons is more
probable
• Structure of electron shell determines chemical
properties
– Valence electrons
– The octet rule
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Periodic table of elements
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Molecular (covalent) compounds
• Consist of particles (molecules) made of several
atoms connected with covalent bonds
• Examples of molecules:
H H
– Gases: diatomic
– H2O, NH3 etc.
H
O
H
H N H
H
– Covalently-bonded crystals: diamond
– Macromolecules of proteins and nucleic acids
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Diamond and graphite
Fig. from Wikipedia
Covalent chemical bond
• Most often between two non-metals
• Chemical bond that results from two nuclei attracting
the same pair(s) of electrons
• Based on electron sharing
H
H
H
H
• If more pairs shared … double or triple bonds
• Some atoms may also have nonbonding (lone) electron
pairs
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Valence
• Number of covalent bonds formed by an
atom
• The octet rule: tendency to achieve electron
configuration of the nearest noble gas
– e.g. H-F: H achieves configuration of He,
F configuration of Ne
• Therefore O usually bivalent, N trivalent,
C tetravalent etc.
Coordination covalent bond
• (also dative, donor-acceptor bond)
• Both bonding electrons provided by one of
the atoms (donor), whereas the other atoms
provides an empty orbital (akceptor)
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Ionic compounds
• Consist of charged
particles (ions) held
together by ionic bonds
• Metal + non-metal(s)
• Cations (+) and anions (-)
combine in space to
achieve electroneutrality
Ions
• Charged because number of electrons does
not match number of protons
• Tendency to form ions depends on
electronegativity of element
• Monoatomic: Na+, Cl-, H+, Fe2+
• Polyatomic: NO3- , SO42• Complex: [Fe(CN)6]4-
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Polyatomic ions of oxo-acids:
e.g. sulfate, SO42- :
resonance stabilization of sulfate ion
..similar is nitrate NO3-, phosphate PO43-, carbonate CO32-, etc.
Coordination (complex) ions
e.g. [Fe(CN)6]4-
• central atom of
transition metal
providing empty
orbitals
• ligands providing free
electron pairs
• Number of ligands
(coordination number)
is usually 4 or 6
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Ionic salts: no true molecule
• Crystal lattice of NaCl:
Formula unit
• Dissolution of NaCl in water: electrolytic dissociation
producing hydrated independent ions Na+, Cl-
Chemical formulas
• Stoichiometric (empirical):
– e.g.: glucose CH2O
• Molecular:
– e.g.: glucose C6H12O6
• Structural:
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Chemical formulas
• Stoichiometric (empirical):
– e.g.: glucose CH2O; sodium chloride NaCl
• Molecular:
– e.g.: glucose C6H12O6; sodium chloride NaCl
• Structural:
Stoichiometric equation:
AgNO3 + KCl
AgCl + KNO3
Ionic equation:
Ag+ + NO3- + K+ + Cl-
AgCl + K+ + NO3-
Net ionic equation:
Ag+ + NO3- + K+ + ClAg+ + Cl-
AgCl + K+ + NO3AgCl
white ppt
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Also possible:
AgNO3(aq) + KCl(aq)
Ag+(aq) + Cl-(aq)
AgCl(s) + KNO3(aq)
AgCl(s)
(aq)
... aqueous
(s)
... solid
(l)
... liquid
(g)
... gaseous
Polarity of chemical bond
Determined by difference in electronegativity
of the two connected atoms:
< 0.4 nonpolar covalent bond
Gradual
e.g. H-H, carbon-hydrogen
transition !
0.4 - 1.7 polar covalent bond
e.g. H-O-H, NH3, carbon-oxygen, carbon-nitrogen
>1.7 ionic bond
e.g NaCl...
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Oxidation number (formal valency)
• Oxidation number of element in compound
equals its charge after giving all bonding
electron pairs to the more electronegative
atom
• Can be zero, positive or negative integer
• Basis for nomenclature of inorganic
compounds
• Redox reactions: oxidation number
increases in oxidation, decreases in
reduction
Rules for determination of oxidation
numbers
• Free electroneutral atom, or atom in molecule of pure
element: oxidation number = 0
• Oxidation number of a monoatomic ion equals its
charge
• In heteroatomic compounds the bonding electrons are
given to the more electronegative atom, practically:
– H has nearly always oxidation number I (only in
metallic hydrides -I)
– O almost always -II (only in peroxides -I)
– F always -I
– Alkali metals (Na, K..) always I
– Alkaline earth elements (Ca, Mg..) always II
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Rules for determination of oxidation
numbers
Examples:
CO2 : CIV, O-II
H2SO4: HI, SVI, O-II
Sum of oxidation numbers of all atoms in electroneutral
molecule is 0, in polyatomic ion equals the ion charge
e.g: CO32-: CIV,
O-II
1×IV + 3 ×(-II) = -2
Czech nomenclature of oxides:
Oxidation number
I
II
III
IV
V
VI
VII
VIII
Suffix
-ný
-natý
-itý
-ičitý
-ečný/-ičný
-ový
-istý
General formula
X2O
XO
X2O3
XO2
X2O5
XO3
X2O7
-ičelý
XO4
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Naming Binary Ionic Compounds
• Compounds of one metallic element and one
nonmetal (e.g. metallic oxides, hydroxides,
halogenides, sulfides)
• Name of metal + stem of nonmetal + -ide
• Examples:
– Al2O3, aluminum oxide
– Ba(OH)2, barium hydroxide
– KCl, potassium chloride
– ZnS, zinc sulfide
Naming Binary Ionic Compounds
• Numerical prefixes are never used.
• If the metal can exist in more oxidation states,
its oxidation number is included to the name
• Examples:
– FeCl3, iron(III) chloride (ferric chloride)
– FeCl2, iron(II) chloride (ferrous chloride)
– CuO, copper(II) oxide (cupric oxide)
– Cu2O, copper(I) oxide (cuprous oxide)
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Naming Binary Molecular Compounds
• Compounds of two nonmetals (e.g. oxides of
nonmetals)
• Name of less electronegative element + stem of
the other element + -ide
• Numerical prefixes precede names of both
nonmetals
• Examples:
– CO, carbon monoxide
– N2O5, dinitrogen pentoxide
– CCl4, carbon tetrachloride
– H2S, hydrogen sulfide
Numerical prefixes:
Number
1
2
3
4
5
6
7
Prefix
monoditritetrapentahexahepta-
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octa-
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nona-
10
deca-
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Naming acids and their salts
Brønsted-Lowry concept of acids and bases:
• Acid is a proton donor
• Base is a proton acceptor
HCl(aq) + H2O(l)
Acid
Base
H3O+(aq) + Cl-(aq)
Conjugate
acid
Conjugate
base
Salt: ionic compound, product of neutralization
reaction between acid and base (the acidic
proton replaced with metal cation)
Naming acids and their salts
A) hydroacids:
Gaseous nonmetallic hydrides whose aqueous
solutions are acidic
E.g. HCl, hydrochloric acid (aqueous
hydrogen chloride), salts: chloride
Likewise:
– HBr, hydrobromic acid, salts bromides
– H2S, hydrosulfuric acid, salts sulfides
– HCN, hydrocyanic acid, salts cyanides
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Naming acids and their salts
B) oxo-acids:
Central atom + -OH groups, protons
dissociate from oxygen. In salts appear as
polyatomic anions
If only one oxidation state of the central atom
is possible:
Stem of the central atom + -ic acid
E.g. H2CO3, carbonic acid, salt: carbonate
Naming acids and their salts
B) oxo-acids:
If there are two possible oxidation states of
the central atom:
Higher ox. number: -ic acid, salt: -ate
Lower ox. number: -ous acid, salt: -ite
Example:
H2SO4, sulfuric acid, salt: sulfate
H2SO3, sulfurous acid, salt: sulfite
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Naming acids and their salts
B) oxo-acids:
If there are more than two possible oxidation
states of the central atom, prefixes are used:
HClO, hypochlorous acid, salt: hypochlorite
HClO2, chlorous acid, salt: chlorite
HClO3, chloric acid, salt: chlorate
HClO4, perchloric acid, salt: perchlorate
The oxidation numbers can also be
used with metals in anions:
MnO42-: manganate(VI) or just manganate
MnO4-: manganate(VII) or permanganate
[Fe(CN)6]4-: hexacyanoferrate(II)
or ferrocyanide
[Fe(CN)6]3-: hexacyanoferrate(III)
or ferricyanide
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